1. INTRODUCTION TO ACID BASE BALANCE
Dr. Sadia Haroon
Lecture # 2

2. Objectives
Explain how the pH of the blood is stabilized by bicarb buffer and define the terms acidosis and alkalosis.
Explain how the acid-base balance of the blood is affected by C02 and HC03-, and describe the roles of the lungs and kidneys in maintaining acid- base balance.
Explain how C02 affects blood pH, and hypoventilation and hyperventilation affect acid- base balance.
Explain how the interaction between plasma K+ and H+ concentrations affects the tubular secretion of these.

3. ACID BASE HOMEOSTASIS
Acid-Base homeostasis involves chemical and physiologic processes responsible for the maintenance of the acidity of body fluids at levels that allow optimal function of the whole individual

4. ACID BASE HOMEOSTASIS
The chemical processes represent the first line of defense to an acid or base load and include the extracellular and intracellular buffers
The physiologic processes modulate acid-base composition by changes in cellular metabolism and by adaptive responses in the excretion of volatile acids by the lungs and fixed acids by the kidneys

5. Acids can be defined as a proton (H+) donor
Hydrogen containing substances which dissociate in solution to release H+
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6. Physiologically important acids include: Carbonic acid (H2CO3)
Phosphoric acid (H3PO4)
Pyruvic acid (C3H4O3)
Lactic acid (C3H6O3)
These acids are dissolved in body fluids
Phosphoric acid
Lactic acid
Pyruvic acid
Carbonic acid

7. Types of Acids in the Body
Volatile acids:
Can leave solution and enter the atmosphere.
H2C03 (carbonic acid).
Pco2 is most important factor in pH of body tissues.

8. Types of Acids in the Body
Fixed Acids:
Acids that do not leave solution.
Sulfuric and phosphoric acid.
Catabolism of amino acids, nucleic acids, and phospholipids.

9. Types of Acids in the Body
Organic Acids:
Byproducts of aerobic metabolism, during anaerobic metabolism and during starvation, diabetes.
Lactic acid, ketones.

10. Bases can be defined as: A proton (H+) acceptor
Molecules capable of accepting a hydrogen ion (OH-)
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11. Physiologically important bases include: Bicarbonate (HCO3- )
Biphosphate (HPO4-2 )
Biphosphate
Bicarbonate

12. H2O H+ + OH- H+ ion is an acid OH- ion is a base pH SCALE
pH refers to Potential Hydrogen
Expresses hydrogen ion concentration in water solutions
Water ionizes to a limited extent to form equal amounts of H+ ions and OH- ions
H2O H+ + OH-
H+ ion is an acid
OH- ion is a base

13. Acid-Base Biochemistry Methods
pH electrode

14. pH = log 1 / H+ concentration
pH SCALE
pH equals the logarithm (log) to the base 10 of the reciprocal of the hydrogen ion (H+) concentration
H+ concentration in extracellular fluid (ECF)
pH = log 1 / H+ concentration
4 X ( )

15. pH = 4 is more acidic than pH = 6
pH SCALE
pH = 4 is more acidic than pH = 6
pH = 4 has 10 times more free H+ concentration than pH = 5 and 100 times more free H+ concentration than pH = 6
ACIDOSIS
NORMAL
ALKALOSIS
DEATH
DEATH
6.8
7.3
7.4
7.5
8.0
Venous Blood
Arterial Blood

16. pH SCALE

17. ACIDOSIS / ALKALOSIS

18. Acidosis Alkalosis ACIDOSIS / ALKALOSIS
A condition in which the blood has too much acid (or too little base), frequently resulting in a decrease in blood pH
Alkalosis
A condition in which the blood has too much base (or too little acid), occasionally resulting in an increase in blood pH

19. CHANGES IN CELL EXCITABILITY
pH decrease (more acidic) depresses the central nervous system
Can lead to loss of consciousness
pH increase (more basic) can cause over-excitability
Tingling sensations, nervousness, muscle twitches

20. INFLUENCES ON ENZYME ACTIVITY
pH increases or decreases can alter the shape of the enzyme rendering it non- functional
Changes in enzyme structure can result in accelerated or depressed metabolic actions within the cell

21. INFLUENCES ON K+ LEVELS
When reabsorbing Na+ from the filtrate of the renal tubules K+ or H+ is secreted (exchanged)
Normally K+ is secreted in much greater amounts than H+ K+ K+ K+ K+ K+ K+
Na+
Na+
Na+
Na+
Na+
Na+ H+ K+

23. ACID-BASE REGULATION

24. Buffer Systems Provide or remove H+ and stabilize the pH.
Include weak acids that can donate H+ and weak bases that can absorb H+.
Does NOT prevent a pH change.

25. Henderson-Hasselbach Equation
HA = weak acid
A- = Conjugate base
1) Ka = [H+][A-]
[HA]
2) [H+] = Ka [HA]
[A-]
3) -log[H+] = -log Ka -log [HA]
[A-]
* H-H equation describes
the relationship between
pH, pKa and buffer concentration
4) -log[H+] = -log Ka +log [A-]
[HA]
5) pH = pKa +log [A-]
[HA]

26. Acid/conjugate base pairs
HA + H2O A- + H3O+
HA A- + H+
HA = acid ( donates H+)(Bronstad Acid)
A- = Conjugate base (accepts H+)(Bronstad Base)
Ka = [H+][A-]
[HA]
Ka & pKa value describe tendency to loose H+
large Ka = stronger acid
small Ka = weaker acid
pKa = - log Ka

28. Henderson-Hasselbach Equation
Consider the dissociation of a general acid HA
HA H+ + A-
We can define a dissociation constant (K) where
Rearranging gives
Taking logarithms on both sides and multiplying by -1 gives:
-log[H+] = -logK – log [HA]/[A-] or
pH = pK + log [A-]/[HA]

29. Henderson-Hasselbach Equation
HA = weak acid
A- = Conjugate base
1) Ka = [H+][A-]
[HA]
2) [H+] = Ka [HA]
[A-]
3) -log[H+] = -log Ka -log [HA]
[A-]
* H-H equation describes
the relationship between
pH, pKa and buffer concentration
4) -log[H+] = -log Ka +log [A-]
[HA]
5) pH = pKa +log [A-]
[HA]

30. Henderson-Hasselbalch Equation
This equation can be used to determine the pH if the pK and ratio of the ionised and unionised forms is known.
The pKa (a for acid) is the –ve log of the dissociation constant of the acid. It is the pH at which the ratio of the ionised and unionised species is equal to 1. ie the molar concentration of the ionised and unionsed species is the same.
Similarly pKb is –ve log of the dissociation constant of the base

31. Regulation of H+ concentration
Concentration of hydrogen ions is regulated sequentially by:
Chemical buffer systems –act within seconds
The respiratory center in the brain stem –acts within 1-3 min
Renal mechanisms –require hours to days to effect pH changes
Sources of hydrogen ions anaerobic and aerobic respiration of glucose incomplete oxidation of fatty acids oxidation of sulfur-containing amino acids hydrolysis of phosphoproteins and nucleic acids

32. Buffers
Definition: A weak acid plus its conjugate base that cause a solution to resist changes in pH when an acid or base are added
Effectiveness of a buffer is determined by:
1) the pH of the solution, buffers work best within 1 pH unit of their pKa
2) the concentration of the buffer; the more present, the greater the buffering capacity

33. The Major Body Buffer Systems
Site
Buffer System
Comment
ISF
Bicarbonate
For metabolic acids
Phosphate
Not important because concentration too low
Protein
Blood
Important for metabolic acids
Haemoglobin
Important for carbon dioxide
Plasma protein
Minor buffer
Concentration too low
ICF
Proteins
Important buffer
Phosphates
Urine
Responsible for most of 'Titratable Acidity'
Ammonia
Important - formation of NH4+
Bone
Ca carbonate
In prolonged metabolic acidosis

35. React very rapidly (less than a second) 2) Respiratory Regulation
ACID-BASE REGULATION
Maintenance of an acceptable pH range in the extracellular fluids is accomplished by three mechanisms:
1) Chemical Buffers
React very rapidly (less than a second)
2) Respiratory Regulation
Reacts rapidly (seconds to minutes)
3) Renal Regulation
Reacts slowly (minutes to hours)

36. Buffering Capacity in Body
52% is in cells, 5% is in RBCs
43% is in the extracellular space
of which 40% by bicarbonate buffer, 1% by proteins and 1% by phosphate buffer system

37. Organs involved in the regulation of A-B-balance
CO2 production from complete oxidation of substrates
20% of the body’s daily production
metabolism of organic acid anions
such as lactate, ketones and amino acids
metabolism of ammonium
conversion of NH4+ to urea in the liver results in an equivalent production of H+
Production of plasma proteins
esp. albumin contributing to the anion gap
Bone inorganic matrix consists of hydroxyapatite crystals (Ca10(PO4)6(OH)2]
bone can take up H+ in exchange for Ca2+, Na+ and K+ (ionic exchange) or release of HCO3-, CO3- or HPO42-

38. Buffer capacity
The buffer capacity of a system is already defined as the amount of strong acid or base added to one litre (l) of the system in order to change the pH one unit

39. Take up H+ or release H+ as conditions change
Control of Acids
Buffer systems
Take up H+ or release H+ as conditions change
Buffer pairs – weak acid and a base
Exchange a strong acid or base for a weak one
Results in a much smaller pH change

40. Take up H+ or release H+ as conditions change
Control of Acids
Buffer systems
Take up H+ or release H+ as conditions change
Buffer pairs – weak acid and a base
Exchange a strong acid or base for a weak one
Results in a much smaller pH change

41. Chemical Buffers Act within fraction of a second. Protein. HCO3-.
Phosphate.

42. Buffer H+ OH- H+ OH- OH- H+ ACID-BASE REGULATION Chemical Buffers
The body uses pH buffers in the blood to guard against sudden changes in acidity
A pH buffer works chemically to minimize changes in the pH of a solution H+
OH- H+
Buffer
OH-
OH- H+

43. Carbonic Acid – Bicarbonate Buffer System
~ Most important in the ECF
pk. = 6.1.
Present in large quantities.
Open system.
Respiratory and renal systems act on this buffer system.

44. Bicarbonate buffer system
Present in intra-and extracellular fluid
Bicarbonate ion acts as weak base, carbonic acid acts as a weak acid
Bicarbonate ions combine with excess hydrogen ions to form carbonic acid
Carbonic acid dissociates to release bicarbonate ions and hydrogen ions
H+ + HCO3-  H2CO3  H+ + HCO3-

45. Bicarbonate buffer
Sodium Bicarbonate (NaHCO3) and carbonic acid (H2CO3)
Maintain a 20:1 ratio : HCO3- : H2CO3
HCl + NaHCO3 ↔ H2CO3 + NaCl
NaOH + H2CO3 ↔ NaHCO3 + H2O

46. HCO3- Limitations Cannot protect ECF from respiratory problems.
Cannot protect ECF from elevated or decreased CO2.
Limited by availability of HCO3-.

48. The blood buffering system, simplified

50. Phosphate Buffer system
~ Important in ICF & urine

51. Phosphate buffer system
Important in intracellular fluid and urine pH regulation
Consists of two phosphate ions
Monohydrogenphosphate ions act as a weak base and combine with hydrogen ions to form dihydrogenphosphate
Dihydrogenphosphate dissociates to release hydrogen ions
H+ + HPO4-2  H2PO4-  H+ + HPO4-2

52. Phosphate & Intracellular Buffers
Both Intra and Extra cellular phosphate act as a buffer . But its role is minor compared to HB or HCO3.
Intracellular buffers are needed because H doesn’t cross Plasma Membrane .
Intracellular PH is more acidic . (7.2)

53. Phosphate buffer Major intracellular buffer H+ + HPO42- ↔ H2PO4-
OH- + H2PO4- ↔ H2O + H2PO42-

55. Phosphate has three ionizable H+ and three pKas

57. Protein Buffer Systems
~ Important in ECF and ICF
~ Interact with other buffer systems

58. Protein buffer system
Consists of Plasma Proteins (albumin, hemoglobin)
Remember proteins are just chains of AAThe exposed amine group of the AA (NH2) accepts H+ ions when conditions are acidic
The exposed carboxyl group of AA can release H+ ions when conditions are basic Proteins can act as Acids or Bases

59. Protein Buffers Includes hemoglobin, work in blood and ISF
Carboxyl group gives up H+
Amino Group accepts H+
Side chains that can buffer H+ are present on 27 amino acids.

60. Hemoglobin is an important blood buffer particularly for buffering CO2
Protein buffers in blood include haemoglobin (150g/l) and plasma proteins (70g/l). Buffering is by the imidazole group of the histidine residues which has a pKa of about 6.8. This is suitable for effective buffering at physiological pH.
Haemoglobin is quantitatively about 6 times more important then the plasma proteins as it is present in about twice the concentration and contains about three times the number of histidine residues per molecule. For example if blood pH changed from 7.5 to 6.5, haemoglobin would buffer 27.5 mmol/l of H+ and total plasma protein buffering would account for only 4.2 mmol/l of H+.

61. The acid-base buffering systems of the body.
The two buffer systems are in dynamic equilibrium with the same hydrogen ion concentration (pH), so that a change induced in the concentration of any one factor in either buffer system rapidly affects the other system and a new hydrogen ion concentration in the blood is established.
The lungs assist in maintaining a constant blood pH by removing CO2,
while the kidney excretes acid in the form of H2PO4- and NH4 and alkali in the form of HCO3-.

62. Chemical Buffers Act within fraction of a second. Protein. HCO3-.
Phosphate.

63. Cell Metabolism Respiratory Regulation ACID-BASE REGULATION CO2 CO2
Carbon dioxide is an important by-product of metabolism and is constantly produced by cells
The blood carries carbon dioxide to the lungs where it is exhaled
Cell Metabolism
CO2
CO2
CO2
CO2
CO2
CO2

64. Respiratory Buffer Systems
The respiratory system regulation of acid- base balance is a physiological buffering system
There is a reversible equilibrium between:
Dissolved carbon dioxide and water
Carbonic acid and the hydrogen and bicarbonate ions
CO2+ H2O ↔H2CO3↔H++ HCO3¯

65. Respiratory Buffer Systems
• CO2 is produced by cellular respiration.
• CO2 is converted to bicarbonate by carbonic anhydrase.
• results in LOWER pH in
respiring tissues.
• CO2 is exhaled in lungs.

66. Haemoglobin binds both CO2 and H+ and so is a powerful buffer
Haemoglobin binds both CO2 and H+ and so is a powerful buffer. Deoxygenated haemoglobin has the strongest affinity for both CO2 and H+; thus, its buffering effect is strongest in the tissues. Little CO2 is produced in red cells and so the CO2 produced by the tissues passes easily into the cell down a concentration gradient. Carbon dioxide then either combines directly with haemoglobin or combines with water to form carbonic acid. The CO2 that binds directly with haemoglobin combines reversibly with terminal amine groups on the haemoglobin molecule to form carbaminohaemoglobin. In the lungs the CO2 is released and passes down its concentration gradient into the alveoli.

67. Respiratory System 2nd line of defense.
Acts within min. maximal in hrs.
H2CO3 produced converted to CO2, and excreted by the lungs.
Alveolar ventilation also increases as pH decreases (rate and depth).
Coarse , CANNOT eliminate fixed acid.

68. Kidney Regulation ACID-BASE REGULATION
Excess acid is excreted by the kidneys, largely in the form of ammonia
The kidneys have some ability to alter the amount of acid or base that is excreted, but this generally takes several days

69. Urinary Buffers Nephron cannot produce a urine pH < 4.5.
IN order to excrete more H+, the acid must be buffered.
H+ secreted into the urine tubule and combines with HPO4-2 or NH3.
HPO4-2 + H H2PO4-2
NH3 + H NH4+